


Acidity and Acidimetry of Soils 



BY 



HENRY GRANGER KNIGHT 

B.A. University of Washington, 1902 
M.A. University of Washington, 1904 



Thesis Submitted in Partial Fulfillment of the Requirements 

for the Degree of Doctor of Philosophy in Chemistry 

in The Graduate School of the University 

of Illinois, 1917 



(Reprinted from the Journal of Industrial and Engineering Chemistry, vol. 12, no. 4, p. 
340; no. S, p. 457; no. 6, p. S59. April, May and June 1920.) 



Acidity and Acidimetry of Soils 



BY 



HENRY GRANGER KNIGHT 

B.A. University of Washington, 1902 
M.A. University of Washington, 1904 



Thesis Submitted in Partial Fulfillment of the Requirements 

for the Degree of Doctor of Philosophy in Chemistry 

in The Graduate School of the University 

of Illinois, 1917 



(Reprinted from the Journal of Industrial and I'.ngineering Chemistry, vol. 12, no. 4, p. 
340: no. 5, p. 457; no. 6, p 5.S9 April, May and June 1920 ) 






LIBRARY OF CONGRESS 

FEB281921 

DOCUMENTS DIVISION 



I Reprinted from the Journal of Industrial and Engineering Chemistry 
Vol. 12, No. 4, page 340. April, 1920.] 



ACIDITY AND ACIDIMETRY OF SOILS. ' I— STUDIES 

OF THE HOPKINS AND PETTIT METHOD FOR 

DETERMINING SOIL ACIDITY 

By Henry G. Knight 

Oklahoma Agricultural and Mechanical College, Stillwater, 
Oklahoma 

Received October 14, 1919 

The Hopkins and Pettit method of determining 
soil acidity^ proposed in 1902 is essentially as follows: 
100 g. of soil are shaken in a bottle of 400 cc. capacity 
with 250 cc. of 5 per cent commercial common salt 
solution for 3 hrs. 125 cc. of the clear liquid are taken 
off, boiled to expel carbon dioxide, and titrated using 
phenolphthalein as an indicator. The results are 
multiplied by 3 as a factor to determine the total 
amount of base required. Later^ a normal solution of 
potassium nitrate was substituted for the 5 per cent 
commercial common salt and the factor 2.5 recom- 
mended. The modified method is still the provisional 
method of the A. O. A. C. for determining the acidity 
of soils. 

Veitch^ criticizes the Hopkins method upon the 
grounds that it indicates only the apparent need for 
lime or the most urgent need, and claims further that 
the acidity shown by this method is largely due to 
aluminates. He also notes that there is a great 
discrepancy between the Hopkins method and that 
proposed by himself^ upon soils high in organic matter. 
Harris" claims that the acidity shown by this method 
is due to selective ion absorption by the soil colloids, 
basing his views upon the fact that the acidity shown 
by the extract is dependent upon the character of the 
salts used. Freer" also holds this view. Truog* 
strenuously combats the theory of colloidal absorption 

1 From a thesis submitted to the faculty of the University of Illinois 
in partial fulfillment of the requirements for the degree of Doctor of Phil- 
osophy. Acknowledgment is made of many helpful suggestions and 
criticisms from Prof C. G. Hopkins and Prof. A. H. Noyes. 

' Nineteenth Annual Proceedings, O. A. C, U. S. Dept. of Agr., Bureau 
of Chemistry, Bulletin 73 (1902), 114. 

3 U. S. Dept. of Agr., Bureau of Chemistry. Bulletin 107 (1908), 20; 
Hopkins, "Soil Fertility and Permanent Agriculture," 1910, 566. 

* J. Am. Chem. Soc, 26 (1914). 637. 

i Ibid., 24 (1902), 1120. 
Michigan Agr. College and Station, Bulletin 19 (1914). 

" Penn. Dept. Agr., Bulletin 261 (1915), 106. 

s J. Phys. Chem., 20 (1916), 157. 

(0 



and brings evidence to support the view of Hopkins 
that the reaction is one of double decomposition be- 
tween the acids or acid salts in the soil and the neutral 
salt solution. Parker^ concluded from analysis of 
extracts prepared by treating soils with potassium 
chloride and potassium acetate that the base was 
absorbed to a little greater extent than it was liberated 
by the soil and that the excess of the anion should be 
accounted for by the presence of the corresponding 
acid. Brogue- states that it has been repeatedly 
proven that the base liberated by the soil is usually 
not merely equivalent to the base absorbed from the 
solution. Sullivan,' Morse and Curry/ Abbott, 
Conn and Smalley,^ Ruprecht,'^ and others have noted 
the presence of aluminum and iron in salt extracts 
from acid soils. Rice^ concludes from hydrogen-ion 
concentration studies upon 31 soils using the indicator 
method of Sorensen* that when so-called acid soils are 
shaken with salt solutions part of the cation of the salt 
is absorbed and an equivalent quantity of the base 
from the soil is given up to the solution. 

It was to test the above points that the following 
investigations were made. 

EXPERIMENTAL 

Harris obtained different lime requirements for soils 
by repeated shaking with different salt solutions. 
These experiments were repeated in this laboratory 
using yellow-gray silt loam, and similar differences were 
obtained as was reported by Harris for different 
salts. As Hopkins claims that the reaction between 
the neutral salt solution and the soil is one of equilib- 
rium, the end reaction would be practically impossible 
to realize by such a treatment. To overcome the 
objections which would arise from the above method 
provisions were made for forcing the salt solutions 
through the soil, so that the soil particles would be 
continually bathed by fresh solutions. 

Twenty grams of yellow-gray silt loam^ were placed 

' This Journal, 6 (1914), 831. 

■ J. Phys. Chem.. 19 (1915), 665. 

s U S. Geol. Survey, Bulletin 312 (1907). 

• New Hampshire Agr. Station, Report 1906-08, 271. 

' Indiana Agr. Expt. Station, Bulletin 170 (1913). 

« Mass. Agr. Expt. Station, Bulletin 161 (1915). 

' J. Phys. Chem.. 20 (1916), 214. 

Biochem. J.. 21 (1909), 131; Walpole, Biochem.. 6 (1911), 207; 8 
(1914), 628. 

» Sample No. 1 . Subsoil from Southern Illinois. Lime requirement : 
Hopkins Method 4.2T., Veitch Method 5.6T., per acre of soil of 2,000,000 
lbs. 

(2) 



upon a dry filter paper and the salt solution was 
allowed to filter through. The filtrate was boiled and 
treated with 0.04 N potassium hydroxide at room 
temperature using phenolphthalein as an indicator, 
with the results shown in Table I. 



Table I — Acidity of Different Fractions of Various Salt Solutions 
Filtered through an Acid Soil 

Equivalent 

Salt . Cc. 0.04 N KOH Required for > to 

Used 100 cc. 250 cc. 250 cc. 250 cc. Total T. CaCOj 

A? KNOb 36.40 2.8 1.2 0.5 40.5 4.05 

A^ KCl 35.95 4.3 2.1 0.5 39.95 3.99 

N NaNOi 26.50 9.2 1.8 1.4 38.9 3.89 

A^ NriCl 31.20 6.4 1.4 ... 39.0 3.90 

A/ CaCh 31.40 5.4 1.3 0.5 38.6 3.86 



The greatest difference shown is 0.19 T., calculated 
as calcium carbonate, which may easily be accounted 
for by errors in reading the end-point. It will be 
noted that the acidity of the sodium nitrate extract 
was quite marked even after 600 cc. had filtered 
through, while the first 100 cc. showed the lowest 
acidity. That none of the extractions were carried to 
completion is evident, but all, with the possible ex- 
ception of that with sodium nitrate, were carried 
to a point beyond which it was impossible to measure 
the acidity with any degree of accuracy by the ordinary 
indicator methods. The calcium salt extract would be 
expected to show a slower reaction after the first 
surface reaction because of the greater insolubility 
of calcium compounds which would be formed upon 
the surface of the soil grains. 

INDICATOR EFFECTS — In the first two series of ex- 
tracts, a precipitate which had a rather marked efifejlt 
upon the indicator was always formed in considerable 
quantity. The pink color produced by the addition 
of a slight excess of base disappeared after a time 
even when the titrating flask was tightly stoppered. 
By the further addition of base the color could be 
brought back. The end-point is also markedly in- 
fluenced by the amount of indicator present. It is 
quite apparent that the indicator is absorbed to a 
marked extent by the precipitate, and instead of the 
simple equilibrium 

HIn -^ H+ + In- 
colorless colored 

there must be taken into consideration the equilib- 
rium with the absorbed indicator 

(3) 



HIn -^^ HIn Z^ H+ + In", 
absorbed colorless colored 

Two equal quantities of a potassium salt extract of 
an acid soil gave readings, as shown in Table II, which 
clearly indicate the variation in results which may be 
obtained by using different amounts of indicator. 

Table II — Effect of Different Amounts of Indicator upon Titration 

End-Point 

Phenolphthalein 0.04 N KOH 

Used Required 

Drops Cc. 

4 62.7 

10 61.1 

In several instances a point was reached which 
showed no visible color change with four or five drops 
of indicator, while upon the addition of larger quanti- 
ties a marked color change was observed. To over- 
come as far as possible the variation due to the indicator 
the same quantity was used in each case unless other- 
wise stated. 

EFFECT OF TEMPERATURE UPON TITRATION The 

temperature at which the titration is carried out was 
found to produce an effect which is shown in Table 
III. Two equal quantities of potassium nitrate 
extract were titrated with 0.04 N potassium hydroxide 
using the same quantity of phenolphthalein as an 
indicator. The only difference between the duplicates 
was that of temperature. 

Table III — Effect of Temperature upon Titration End- Point 
Temp. " C. Cc. KOH 

22 16.7 

85 19.2 

EFFECT OF TEMPERATURE UPON THE AMOUNT OF 

ACIDITY SHOWN — If the Hopkins method is a measure 
of the colloidal adsorption for the base by a soil as 
maintained by Harris^ there should be a temperature 
effect which could be measured. Travers'^ has shown 
that the adsorption of carbon dioxide by charcoal 
decreases markedly with rise in temperature, and, we 
may expect a similar change to be shown by soils in 
contact with neutral salt solutions. 

To test this theory an apparatus, shown in Fig. i, 
was arranged in a constant temperature electrical oven 
of the Freas type. The apparatus was arranged in 
duplicate, a is the receptacle for the neutral salt 
solution, b a stopcock for regulating the flow of the 
neutral salt through tube d which passes through the 
ventilating openings c provided by the manufacturers 

' Loc. cil. 

2 Proc. Roy. Soc. London, 74 (1904), 126. 

(4) 



in the stock oven. The bulb e serves to bring the salt 
solution to the temperature of the oven before it 
runs into the receptacle / which contains the soil under 
investigation. The filtrate passes out of the ap- 
paratus through tube g, through ventilating openings 
in the oven at c' , and is caught in measuring flasks; 






Fig. 1 

i is a thermometer placed in the center of the oven for 
noting the temperature. 

Twenty grams of soil were placed in the receptacle, 
normal potassium nitrate solution allowed to percolate 
through, and titrated at room temperature with results 
as shown in Table IV. 

The temperature eflfect is very slight, and certainly 
does not indicate colloidal adsorption, unless this may 
be an exceptional case. 

(5) 



Tabi.u IV' — Effect of Temperature upon Reaction between Acid 
Soil and Neutral Salt Solution 

■ Temperature . 

90° + 1° -. 25° ± 1° . 

Pbrcolati- 0.04 N KOH CaCOa 0.04 7^ KOH CaCOa 

Cc. Cc. T. Cc. T. 

100 37.85 3.785 36.4 3.64 

400 3.15 0.315 3.0 0.30 

Total 500 41.00 4.1 39.4 3.94 

EFFECT OF STRENGTH OF SALT SOLUTION PotaSSium 

nitrate solut'Ions of different strengths were filtered 
through 2o g. samples of yellow-gray silt loam and the 
filtrate titrated with 0.04 N KOH, using phenol- 
phthalein as an indicator with results shown in Table V. 

Table V — Ivffect op Strength of Salt Solution upon the Total 
Amount of Acidity 

Fractions • Cc. 0.04 A' KOH Required 

Cc. N KNO3 0.5 N KNOj 0.1 N KNOi 

1 — 100 35.0 33.3 18.3 

2—100 3.5 4.7 7.6 

3—100 1.2 1.5 4.3 

4—100 0.8 0.9 3.7 

5 — 100 0.5 0.6 3.3 

6—100 .. .. 1.6 

Total 40.8 41.0 38.8 

asT. CaCOi 4.08 4.10 3.88 

It will be noted that with the stronger salt solutions 
the greatest acidity is shown in the first loo cc. but 
rapidly falls while the weaker solutions show a higher 
acidity in succeeding fractions. Evidently end ex- 
tractions would lead to the same end results regardless 
of the strength of the salt solution. 

Further study of this reaction was made by ex- 
tracting 10 g. samples of the same soil with various 
strengths of potassium nitrate solution and titrating 
the acid in the first loo cc. of the percolate. The 
results are tabulated in Table VI. 

Table VI — Acidity Shown in First 100 Cc. Filtrate from Acid Soil, 

Using Neutral Salt Solutions of Various Strengths 

0.04 N KOH 

Normality of Required Calculated as 

KNO3 Cc. T. CaCOa 

1.00 15.2 3.40 

0.50 15.3 3.06 

0.25 15.2 3.04 

0.125 13.25 2.65 

0.0625 8.00 1.60 

0.0312 4.10 0.82 

0.0156 1.95 0.39 

0.0078 1.17 0.23 

The acidity of the first portion of the extract in- 
creases with increase in concentration of the neutral 
salt solution. 

Table VII — Amounts op Lime Absorbed by Acid Soil from Solutions 
OF Various Strength 



Lime Added 


Absorbed 


Left in 


Final 


Calculated as 


by Soil 


Solution 


Normality of 


T. CaCOj 


T. CaCO. 


T. CaCO. 


Solution 


12 


11.71 


0.29 


0.000285 


14 


13.56 


0.64 


0.00064 


16 


14.59 


1.41 


0.0014 


20 


16.75 


3.25 


0.0032 


30 


20.93 


9.07 


0.009 


40 


24 . 96 


15.04 


0.015 



(6) 



EXAMPLES OF SO-CALLED COLLOIDAL ABSORPTION 

As examples of what are usually considered colloidal 
phenomena the following experiments are submitted. 

EXPERIMENT I — 20 g. of yellow-gray silt loam were 
shaken with various amounts of lime contained in 
200 cc. of solution for 1 2 hrs. and aliquot parts of the 
clear liquid were titrated with results as shown in 
Table VII. 

A greater portion of the lime is absorbed from the 
dilute solutions than from the more concentrated, 
thus apparently following the colloidal absorption 
law. 

EXPERIMENT II — 20 g. of soil were 
placed in an extracting apparatus 
{b, Fig. 2), and a 0.04 N calcium 
hydroxide solution allowed to per- 
colate through. In the diagram a is 
the receptacle for the base, e a stop- 
cock for regulating the flow of the 
base into b, and c a graduated re- 
ceiver connected with tube d which 
serves to equalize the pressure. 
The apparatus is a closed system 
and duplicate results were easily 
obtained. 237.2 cc. of filtrate^ 
passed through before a pink color 
could be detected, with phenolphthal- 
ein as an indicator, representing 
a lime adsorption at this point of 
23.7 T. as calcium carbonate. At 
the close of the experiment 537.6 
cc. of filtrate had percolated through, 
the last 50 cc. being 0.0285 N base, 
while the soil had absorbed a total of 
35.28 T. of lime as calcium carbon- 
ate. 15.82 T. of lime were washed out 
by the first 700 cc. of distilled water, 
the end fraction passing through 
0.00242 N alkali. 19.46 tons of 
lime were still left in the soil. 

The details are given in Table 
VIII, in which A and B are duplicate 
determinations. 




Fig. 2 



> Figures given are average of results A and B below. 



(7) 



Table VIII — Amount op Lime Absorbed by Acid vSoiu prom Solution op 

Constant Strength Percolating through the Soil 

A 

Ca(OH)2 





Calculated as 




Calculated as 


Percolate 


T. CaCO. 


Per cent 


T. CaC03 in 


Cc. 


Absorbed 


Absorbed 


Percolate 


236.5 


23.65 


100.0 


0.00 


27.0 


2.50 


92.6 


0.20 


28.5 


1.95 


68.4 


0.90 


45.0 


1.97 


43.7 


2.53 


51.4 


1.61 


31.3 


3.53 


48.6 


0.93 


19.1 


3.93 


50.0 


1 .04 


20.8 


3.96 


50.5 


1.45 


28.7 


3.60 


1. 537.5 


35.10 

B 






238.0 


23 . 80 


100.0 


0.00 


43.2 


4. 17 


96.6 


0.15 


58.0 


3.04 


52.2 


2.76 


50.5 


1.24 


24.5 


3.81 


50.0 


1.08 


20.2 


3.92 


50.0 


1.15 


23.0 


3.85 


50.0 


0.99 


19.8 


4.01 


L 539.7 


35.47 







At the completion of the experiment water was 
added and the first 700 cc. carried through the equiva- 
lent of 16 tons of lime as CaCOs, or nearly half the 
amount absorbed. 

EXPERIMENT III — Twenty gram samples of yellow-gray 
silt loam were shaken for 3 hrs. with 200 cc. of 0.04 N 
calcium hydroxide and potassium hydroxide, respec- 
tively. Since the solution containing the potassium 
hydroxide would not settle, 25 cc. of normal potassium 
nitrate were added to both the calcium and potassium 
hydroxide solutions, and filtered. Titrations of 100 cc. 
of filtrate with 0.04 A^ hydrochloric acid gave the 
results included in Table IX. 

Tablb IX — Amount of Base Not Absorbed by an Acid Soil from Solu- 
tions of Equivalent Strength 

0.04 N HCl Required to 
Neutralize Filtrate 
Base Usbd Cc. 

Ca(OH)2 7.7 

KOH 31.9 

As the potassium hydroxide solution was neutralized, 
a copious precipitate of aluminum hydroxide was 
formed. No precipitate was noted upon neutralizing 
the calcium hydroxide solution. 

It would appear from this experiment that the soil 
has a greater absorption power for calcium than for 
potassium, which is not indicated by other experi- 
ments. A chemical difference in the action of the 
two bases seems the more simple explanation. Potas- 
sium aluminate is soluble while calcium aluminate is 
not. Both are unstable except in the presence of a 
base. Since the potassium aluminate passes into 
solution it is titrated above while the calcium aluminate 
is precipitated around the soil particles. 

(8) 



This apparently throws doubt upon the magnitude 
of the colloidal adsorption effects which may be 
assumed from Expts. I and II above. The probable 
cause would seem to be precipitation effects. Hy- 
drolysis will account for the rapid washing out of the 
lime when water is added in Expt. II. 

EXCHANGE OF BASE 

There is considerable doubt whether, when an acid 
soil is shaken up with a neutral salt, there is a complete 
exchange of base. Rice^ comes to the conclusion that 
there is an equivalent exchange and that the acidity 
is due to the aluminum salts. He further claims that 
the ordinary methods of analysis are too crude to 
determine this accurately. Sharp and Hoagland^ by 
use of the hydrogen electrode show that soil acidity 
is due to an excess of hydrogen ions, and that the 
acidity is increased by the presence of certain neutral 
salts. 

Analysis was made of the potassium nitrate extract 
of yellow-gray silt loam with results shown in Table 
X. 

Table X — Analysis of Potassium Nitrate Extract of an Acid Soil 

627 Cc. of 0.04 N Acid = 
Gram 

Si02 0.04741 

P2O5 0.00576 

AI2O3 0.38822 

Fe203 Trace 

CaO 0.06025 

MnsOi 0.04839 

MgO 0.19901 

The acid combined with the alumina would be 
equivalent to 570 cc. of 0.04 N acid leaving an ex- 
cess of 57 cc. to be accounted for in other ways. 

Since, as has been shown by Blum,^ alumina is com- 
pletely precipitated before the hydrogen-ion concentra- 
tion drops to a value of io~^ and, conversely, alumina 
will not pass into solution until *the hydrogen-ion 
concentration reaches a value higher than io~^, it is 
reasonable to conclude that there must be some ab- 
sorption of base before the alumina will pass into 
solution. If this is true we must assume that the 
dissolution of alumina is a secondary reaction, and the 
above analysis certainly points in this direction. 

It was noted that strongly acid extracts of this 
soil were highly colored with iron. To test the question 
of why iron is not also taken out in larger quantities 

' Loc. cii. 

« J. Agr. Res., [3] 7 (1916). 124. 

' J. Am. Chem. Soc, 7 (1916), 1282. 

(9) 



by a similar secondary reaction, loo g. samples of 
yellow-gray silt loam were shaken with 250 cc. of 
0.04 N acids and a partial analysis made of 125 cc. 
of the filtered extracts. The results are given in 
Table XI. 

Table XI— Partial Analysis op Weakly Acid Extracts op an 
Acid Soil 

AbOj 

.\ciD Used Gram FejOs 

Acetic . 00383 Trace 

Nitric . 069 1 3 Trace 

Hydrochloric 0.07049 Trace 

The solutions were still acid and upon neutralization 
white precipitates of aluminum hydroxide were formed 
in the hydrochloric and nitric acid extracts but none 
was noted in the acetic acid extract until it was 
neutralized and boiled. It would seem that when 
dilute acids are allowed to act upon an acid soil alumina 
is first brought into solution, and that neutral salts, 
when brought into contact with an acid soil, show 
the properties of a weak acid in this respect. 

DIALYSIS 

Normal potassium nitrate solution was shaken with 
yellow-gray silt loam and allowed to settle. The 
supernatant liquid was drawn off into a collodion 
bottle and subjected to dialysis with the results given 
in Table XII. 

Tablu Xli— Dialysis of a Neutral Salt in Contact with an 
Acid Soil 

0.04 A' KOH to Neutralize 
Cc. 

1 St water fraction 3,S . 1 

2nd water fraction 1 2 . 06 

3rd water fraction 24.9 

Left in flask 24.9 

Total 96.96 

74 per cent of the titratable acid had passed through 
the membrane. As the liquid left in the dialyzing 
flask was being titrated a heavy precipitate formed, 
while that which passed through remained clear upon 
neutralization. It was evident that the acid passed 
through while the aluminum hydroxide did not. 
This, however, was to be expected, as this is one of the 
recognized methods for the preparation of colloidal 
aluminum hydroxide.' 

DISTILLATION OF SOIL EXTRACT 

Attempts were made to remove acid from potas- 
sium nitrate and chloride extracts of soil by prolonged 
distillation with steam but without success. Better 
success followed the distillation of the potassium 

> Graham, Ann.. 121 (1866). 41. 

fio) 



acetate extract of the yellow-gray silt loam as shown in 
Table XIII. 

Table XIII — Distillation of a Potassium Acetate Extract of an 
Acid Soil, and of Stock Solution op Potassium Acetate 
(Acidity in Terms of 0.04 N Base) 
Potassium Acetate Potassium Acetate 

Extract ^ . Stock Solution > 

Distillate Residue Distillate Residue 

72.9 Cc. 21.9 Cc. 8.6 Cc. 1.95 Cc. 

Phenolphthalein was used as an indicator. This 
experiment shows the presence of appreciable quanti- 
ties of acetic acid in the soil extract. About three- 
fourths of the acid shown by the extract was distilled 
over. 

Walter Crum^ prepared colloidal aluminum hy- 
droxide by separating the acetic acid by heating, but 
as there were only traces of aluminum salts carried 
by the potassium acetate extract^ it can hardly be 
conceived that the phenomena may be accounted for 
by the presence of these salts, but rather that there is 
an excess of acid. 

COMPARISON OF CATION AND ANION ABSORPTION 

To compare the cation and anion absorption of the 
yellow-gray silt loam from neutral salt solution a 
0.0358 N solution of calcium chloride was allowed 
to percolate through 20 g. of the soil in the apparatus 
shown in Fig. 2. The extract was analyzed for calcium 
and chlorine. The calcium was determined in an 
aliquot portion of the extract which had been freed 
from iron and aluminum by first precipitating as the 
oxalate and titrating the precipitate with a standard 
potassium permanganate in the presence of dilute 
sulfuric acid. The chlorine was determined by the 
Volhard method^ using 0.04 N solutions. The re- 
sults are tabulated in Table XIV. 

Table XIV — Analysis of Fractions op Extract op an Acid Soil bv a 
Neutral Salt Solution to Determine Cation and Anion Absorption 
Extract 
Fractions Calcium Chlorine Acid 

Cc. Normality Normality Normality 

1—50 0.0180 0.0348 0.0046 

2 — 50 0.0314 0.0358 0.0037 

3 — 50 0.0332 0.0358 0.0028 

4 — 50 0.0335 0.0358 0.0022 

The cation is absorbed to a measurable extent, for 
which the change in acidity fails to account. There 
must, therefore, have been an exchange of base. 
This confirms the fact that there is a basic exchange 
regardless of the neutral salt used. (See analysis of 

i Ann., 89 (1854), 168. 

» Compare Conner, Tras Journal, 8 (1916). 35. 

* Ann., 190, 1. 



potassium nitrate extract.) The anion is absorbed 
little or not at all. The slight absorption shown in the 
first so CO. of extract is probably due to the wetting 
of the particles and to a slight dilution of the extract 
by moisture in the air-dried soil. A small amount 
of precipitate was formed in each case upon neutraliz- 
ing the extract. 

SUMMARY 

I — When normal solutions of potassium nitrate, 
potassium chloride, sodium nitrate, sodium chloride, 
and calcium chloride were percolated through an acid 
soil all gave the same end titrations, using phenol- 
phthalein as an indicator. This.corroborates Hopkins' 
statements. 

2 — The acidity of the salt extract of an acid soil is 
independent of the temperature within the range 
from 25° to 90° C. 

3 — The precipitate formed in titrating the soil 
extract obtained by the Hopkins method absorbs the 
indicator to a marked extent. The end result depends 
upon the temperature, time, and amount of indicator 
used. 

4 — The acidity of the first portions of the neutral 
salt extracts of an acid soil increases with increase in 
concentration of the neutral salts. 

5 — The difference in absorption of calcium and 
potassium from solutions of their bases by an acid soil 
may be accounted for by precipitation effects. 

6 — There is a marked basic exchange when a neutral 
salt solution is added to an acid soil, by which alumina 
is carried into solution. This, however, does not 
account for the total acidity of the solution. 

7 — When acid soil is extracted with potassium 
acetate solution, a portion of acetic acid may be 
distilled off from the extract, showing the presence 
of free acid. 

8 — Exchange of acid radicals when an acid soil is 
treated with a neutral salt solution was not noted. 



(12) 



[Reprinted from the Journal of Industrial and Engineering Ciiemistry, 
Vol. 12, No. .5, page 457. May, 1920.] 



ACIDITY AND ACmiMETRY OF SOILS. ^ H— INVESTIGA- 
TION OF ACID SOILS BY MEANS OF THE 
HYDROGEN ELECTRODE 

By Henry G. Knight 

Oklahoma Agricultural and Mechanical College, Stillwater, 

Oklahoma 

Received October 14, 1919 

INTRODUCTION 

Although the literature upon the subject of soil 
acidity is voluminous, the use of the hydrogen elec- 
trode in soil investigations has been rather limited. 

Gillespie- made use of the hydrogen electrode for 
determining the hydrogen ion concentration of a 
mixture of soil with pure water; but, for reasons which 
will develop in the investigations given herewith, the 
presence of a conducting medium was found to be 
desirable. Sharp and Hoagland^ studied the hydrogen 
concentrations of suspensions of soils in pure water 
under various conditions, effects of natural salts and 
bases upon hydrogen ion concentrations of soil sus- 
pensions, and made titrations with various bases. 

The purpose of the present investigation is to study: 
(i) the speed of reactions between neutral salt solutions 
and soils; (2) the speed of reactions in the presence of 
a base; (3) the change in hydrogen ion concentration 
with change of amount of base and with time, and the 
change in conductivity of soil solutions. 

APPARATUS 

Preliminary experiments were conducted with an 
apparatus similar to that described by Hildebrand,'' 
and some experiments described elsewhere with soil 
solutions were carried out, but for use with solutions 
in contact with the soil it was found to be rather unsat- 
isfactory. For all work reported, unless otherwise 
stated, a high-grade potentiometer (Leeds and North- 
rup. No. 28952) was used with a gas cell especially 
designed for the work. 

GAS CELL — Preliminary experiments showed that for 

> This is a thesis submitted to the faculty of the University of Illinois 
in partial fulfillment of the requirements for the degree of Doctor of 
Philosophy. Acknowledgment is made of many helpful suggestions and 
criticisms received from Prof. C. G. Hopkins and H. A. Noyes. 

2 J. Wash. Acad. Set., 6 (1916), 7. 

3 J. Agr. Res., 7 (1916), 124. 

* J. Am. Chem. Soc, 36 (1913), 847. 

(l) 



uniform results it was necessary to have a gas cell 
which could be agitated continuously, as apparently 
the agitation produced by the entering gas was not 
sufficient. After a number of trials the gas cell shown 
in Fig. I was designed for this work. It is cylindrical 
in shape, 3.3 cm. in diameter and 16 cm. in length, 




Fig. I 



the ends being rounded off. At one end an opening 
is provided of a size to carry a No. 4 rubber stopper, 
through which pass the electrical connection a to the 
platinum plate b to serve as the hydrogen electrode, 
the tube c for the ingress of hydrogen gas, and tube d 
for the outlet. Tubes c and d have capillary tubes 
sealed into the ends to regulate the flow of hydrogen. 
To make connections with the calomel half cell a glass 
tube, g, is provided at the further end of the gas cell 
provided with a stopcock, /, and a constricted tip, h. 
This tube reaches to within a few mm. of the bottom 
of the gas cell and as it did not readily clog with soil 
was found to be very satisfactory. The gas cell was 
designed to be of 100 cc. capacity and to be filled half 
full of liquid, leaving room for 50 cc. of gas. The 
reason for this arrangement will develop later. 

The hydrogen electrode 6 is a rectangular piece of 
sheet platinum, 1.2 x 2.4 cm., with pieces of platinum 
wire welded to each end, and is similar to that used by 
Gillespie.' The wires were welded into the glass tube 
at each extremity as shown, making connections with 
the mercury in the tube and at the same time support- 
ing the electrode rigidly. 

Fig. 2 shows diagrammatically the method of agitating 
the cell and connections as measurements were being 
made, a is the calomel electrode provided with a stop- 

< Loc cit 

(2) 




(3) 



cock, b, and c is a U-tube filled with normal potassium 
chloride, making connection with the hydrogen gas 
cell by the tube d. The apparatus was placed upon a 
tilting table hinged at the point o to the base P. The 
table is rocked by means of the adjustable arm 5 and 
crank j attached by suitable gears and pulleys to a 
constant source of power. In building up the shaking 
apparatus free use was made of the parts of one of the 
popular metal construction shapes to teach children to 
make their own toys. 

It was found by experiment that by giving the table 
72 complete oscillations per minute through a total 
angle of 6 ° to 7 ° the liquid in the gas cell was thoroughl y 
agitated; that the rubber stopper was not wetted, but 
that the soil was kept thoroughly mixed with the 
solution while the coarser particles of sand in the soil 
would tend to collect in a nodal point n in the bottom 
of the cell near the center. At this speed the hydrogen 
bubbles rose to about the point m before they broke, 
causing the cell to be quickly freed of air. Except for 
very light peat soils there was no tendency for the 
soil to ascend into the tube d. 

The arrangement of the different parts of the appara- 
tus as used in making measurements is shown diagram- 
matically in Fig. 3. The apparatus is designed to carry 
two gas cells, marked E. M. F., so that duplicate mea- 
surements may be made without changing cells. 

The use of the U-tube for making connections 
between the gas cells and calomel electrode was found 
to be a most satisfactory arrangement. The resistance 
of the chain is slightly increased, but diffusion is 
markedly reduced. During several weeks of con- 
tinuous use, the calomel electrode changed less than 
0.0005 volt. The gas cell is claimed to have the follow- 
ing advantages: 

(i) large capacity, (2) absence of dead airspace. (3) ease of clean- 
ing, (4) all-glass contact with liquid, (5) freedom from clogging, 
(6) ease of manipulation, (7) adaptability to thorough agitation, 
(8) minimum possibility of diffusion. 

CALOMEL ELECTRODE — Calomel was prepared by 
treating pure mercury with dilute nitric acid, precipi- 
tating with hydrochloric acid, washing twenty times 
with potassium chloride, and finally shaking with 
normal potassium chloride and pure mercury, as 
recommended by Ellis^ for obtaining an electrode of 
constant potential. The electrodes were prepared 
with this gray mixture of mercury, calomel, and normal 

J. Am. Chem. Soc, 38 (1916). 737 
(4) 



potassium chloride by first putting into the bottom of 
the carefully cleaned electrode cell a small amount of 
pure mercury to cover the platinum connection. No 
difficulty was experienced in getting electrodes to check 
within 0.0005 volt. The stopcock upon the tube leading 
from the calomel electrode was well greased and was 
opened only when readings were being taken. Con- 
nection was made to the U-tube c (Fig. 2) through a 
rubber stopper well driven home, making this side of 
the system practically gastight. 

HYDROGEN ELECTRODE — The platinum electrode was 
prepared by plating as a cathode in a one per cent 
solution of platinum chloride, containing a small 
amount (about 0.05 per cent) of lead acetate to cause 
the platinum black to adhere. Attempts were made 
to use a solution of pure platinic chloride as recom- 
mended by Ellis, but this proved unsatisfactory as the 
platinum black invariably washed off within a few 
minutes after being placed in the cell. 

AceuhiuLKreR 




r€^ 



+ + - + - 

POTEMT IOMETER 



Fig. 3 

In plating a platinum anode was used. The strength 
of the current was varied from time to time but no 
variation in the potentials of the electrode or in the 
time required to become saturated was noted from this 
cause. If the evolution of hydrogen was too lively 
some of the platinum black was loosened. The plating 
was continued from i to 2 hrs., or until a good heavy 

•'5) 



deposit of platinum black was formed. The electrodes 
were placed in distilled water and given a final wash 
just before using. 

HYDROGEN GENERATOR — Hydrogen was prepared 
electrolytically as needed, using potassium hydroxide 
and nickel electrodes, as shown in Fig. 4. Two salt- 
mouth bottles, G and G', were used. E and E' are the 
nickel electrodes, M and M' are glass partitions run- 
ning almost to the bottom of the cells to reduce diffu- 
sion as much as possible. P and P' are safety valves 
made of test-tubes. The outlets and 0' were closed 
while hydrogen was being taken off. By using two 
cells in parallel and electrolyzing with a current of 10 
amperes, about 150 cc. of hydrogen per minute were 
obtained. The usual general precautions were taken 
in washing and purifying the gas. 

PROCEDURE 

To fill the tube g of the gas cell (Fig. i) the cell was 
partly filled with liquid, a stopper inserted in the 
opening e, and pressure exerted, so that, when the 
stopcock / was opened, the liquid would fill the tube g 
and flow out at h. Care was taken to thoroughly wet 
the stopcock / by loosening and turning. All measure- 
ments were made with this stopcock closed. 

After the gas cell was filled with soil and solution 
and placed as shown in Fig. 2 the table was oscillated 
and hydrogen run in at / for 4 or 5 min.; the tube g 
was closed and rocking continued for 4 min.; hydrogen 
was again run in for 4 min. to drive out the last traces 
of air; g was again closed, the stopcock b opened, and 
readings were taken immediately. 

The procedure varied somewhat with the number 
of readings to be taken, type of experiment, etc. If 
readings were to cover a period of half an hour or more, 
after the first large volume of hydrogen was run in, 
it was found advisable to allow a small amount to 
bubble through to overcome any diffusion which 
might take place through the rubber connections. 

DESCRIPTION OF SOILS INVESTIGATED 

Yellow-Gray Silt Loam, Sample No. i — Anacidsub soil 
collected from the southern part of the state of Illinois. 
This soil gave a lime requirement by the Hopkins 
method* of 4.2 tons and by the Veitch method^ 5.6 

> U. S. Dept. of Agr.. Bureau of Chem., Buttetin 107 Revised (1908), 20- 
* J. Am. Chem Soc, 24 (1902). 1120. 

(6) 



tons calculated as calcium carbonate.' No carbonates 
were present as determined by the Marr- method. 

Black Peaty Loam, Sample No. 1-281^- — -Collected by 
C. G. Hopkins near Bolton, N. C. Limestone require- 
ment by the Hopkins method was 3.445 tons. 

Black Clay Loam, Sample No. 1-284 — Collected by 
C. G. Hopkins near Bolton and Byrdsville. Some leaf 
mold was present. Limestone requirement by the 
Hopkins method was 4.982 tons. 

Peat, Sample No. 1-241 — Labeled deep peat or 
muck, collected near Titusville, Florida, by C. G. 
Hopkins. Limestone required by the Hopkins method 
was 3.267 tons. 

Black Muck, Sample No. 1-242 — ^CoUected from 
Wauchula, Florida, collected by C. G. Hopkins. Lime- 
stone requirement by the Hopkins method was 4.056 
tons. 

h 




Yellow Silt Clay, Sample No. 2660 — Collected from 
Clay County, Illinois. A silty clay, stiflF and plastic. 
Subsoil sample. Limestone required by the Hopkins 
method was 5.08 tons. 

Gray Clayey Silt, Saftiple No. 2g68 — Collected from 
Jackson County, Illinois. A compact, impervious 
subsoil containing iron blotches and concretions. 
Limestone required by the Hopkins method was 7.92 
tons. 

' All figures are based upon an acre of soil 6^/3 in. in depth, calculated 
to weigh 2,000,000 lbs., except peats which are calculated to weigh 1,000,000 
lbs. 

2 J. Agr. Sci., [II], 3, 155. 

3 Numbers refer to the University of Illinois Soil Survey Numbers. 

(7^ 



Gray Plastic Clay, Sample No. 3057 — Collected from 
Winebago County, Illinois. A subsoil sample con- 
taining clayey sand. Limestone required by the Hop- 
kins method was 2.6 tons. 

Yellow Plastic Clayey Silt, Sample No. 3556 — Col- 
lected from Clay County, Illinois. A subsoil sample. 
Limestone required by the Hopkins method was 4.67 
tons. 

Yellow Silt Loam, Sample No. 4068 — Collected from 
La Salle County, Illinois. This was subsurface sample 
having a lime requirement by the Hopkins method of 
2.89 tons. 

Brown Sandy Loam, Sample No. 6316 — This was 
subsurface sample having a limestone requirement by 
the Hopkins method of 1.65 tons. 

EXPERIMENTAL 
EXPERIMENT I. SPEED OF REACTION BETWEEN AN 

ACID SOIL AND A NEUTRAL SALT — This experiment was 
planned to observe the speed of change in hydrogen 
ion concentration of a neutral salt solution when 
shaken together with an acid soil. 

The gas cell was filled with 50 cc. of 0.5 N salt solu- 
tion and after the potential had become constant, 
which was usually within 10 min., 5 g. of soil (2.5 g. 
in the case of peat soil) were quickly introduced and 
readings were taken at stated intervals. 

Table I — Speed of Reaction between Acid Soil and a Neutral Salt 
Solution 





(Readings 


are given 


in volts. Temperature 25° C.) 












Brown 


Yellow 












Sandy 


Plastic 






Yellow-Gray Silt Loam . 


Loam 


Clayey Silt 


Peat 


Time 


0.5 JV 


'A Mol. 


0.5 A' 


0.5 .V 


0.5 N 


0.5 AT 


Min. 


KCl 


CaCh 


KC2H3O2 


KCl 


KCl 


KCl 





0.6969 


0.6823 


0.7295 


0.6966 


0.6967 


0.6967 


5 


0.4843 


0.4947 


0.6715 


0.5658 


0.4944 


0.5706 


10 


0.4843 


0.4950 


0.6707 


0.5660 


0.4944 


0.5691 


15 


0.4848 


0.4950 


0.6706 


0.5660 


0.4944 


0.5687 


20 


0.4854 


0.4950 






0.4944 


0.5687- 


25 


0.4860 




'. 6693 




0.4946 


0.5688 


30 


0.4865 


0.4950 




0^5660 


0.4946 


0.5690 


40 


0.4871 




0^6693 




0.4950 


0.5692 


60 


0.4878 


0^4957 


0.6693 


o!5664 




0.5696 


90 


0.4890 


0.4963 










120 








0^5670 


0^4969 




180 


0^4895 


0.4972 


0.6688 




0.4984 




240 


o!4963' 






o!57332 


0.5010 
0.5063S 




' After 24 hrs. = 


After 48 h 


rs ' After 


12 hrs. 







Under the conditions of the experiment a voltage of 
0.6967 represents a hydrogen ion concentration of 
io~^ or true neutrality. Voltages less than this figure 
show a higher hydrogen ion concentration or a condi- 
tion of acidity while figures above show a hydrogen 
ion concentration of less than io~' or an alkaline 
condition. 

(8) 



In Table I it will be noted that the lowest readings 
were obtained within 5 min. in all cases but two, the 
yellow-gray silt loam with potassium acetate, and 
peat with potassium chloride. The potassium acetate 
solution showed a marked alkalinity before the soil 
was added, therefore we should expect it to act more as 
a base. It will be discussed in connection with the 
next experiment. The peat did not wet readily which 
may account for its behavior. After the low readings 
were obtained there was a gradual increase in the 
voltages, corresponding to a decrease in the hydrogen 
ion concentration. 

It is evident that the main reaction between the 
salt solution and the soil reaches an equilibrium very 
quickly. Secondary reactions are indicated by an 
increase in voltage after a lapse of time, except in the 
case where potassium acetate was used, where secondary 
reactions are not apparent to any great extent. Sec- 
ondary reactions of greater or less magnitude should 
be expected when the hydrogen ion concentration 
rises appreciably above io~^. Reduction of nitrates^ 
may be put forward as an explanation but this is 
immediately questioned since there is no evidence of 
the potassium acetate mixture showing a change 
toward alkalinity. 

EXPERIMENT II. SPEED OF REACTION IN THE PRES- 
ENCE OF A BASE — This experiment was similar to 
Expt. I. 50 cc. of 0.5 N potassium chloride solution 
containing given quantities of calcium hydroxide, cal- 
culated as calcium carbonate, were used, as indicated 
in Table II. Five grams of soil were «sed except in 
the case of peat, when 2.5 g. were used. 

Table II — Speed of Reaction between Soil and Salt Solution in 
Presence of Base 

Brown Yellow 

Sandy Yellow-Gray Plastic 

Loam Silt Loam Clayey Silt Peat 

Time 2 T. 4 T. 10 T. 5 T. 10 T. 

Min. CaC03 CaCOa CaCOa CaCOs CaCOs 

0.9543 0.9715 0.9961 0.9803 0.9961 

5 0.7783 0.7483 0.8505 0.6562 0.8345 

10 0.7565 0.6923 0.8410 0.6302 0.7642 

15 0.7440 0.6467 0.8363 0.6228 0.7370 

20 0.7313 0.6161 0.8356 0.6199 0.7173 

25 0.7269 0.6037 0.8337 0.6170 0.7053 

30 0.7232 0.5978 0.8326 0.6148 0.6951 

40 0.7165 0.5935 0.8298 0.6111 0.6809 

60 0.7113 0.5897 0.8273 0.6072 0.6672 

90 0.5855 0.8262 .... 

120 0.7066 .... 0.5987 0.6499 

180 . 0.5820 0.8206 0.5951 0.6432 

0.7022' 0.81862 0.5933^ 0.6394' 

1 After 48 hrs. ^ After 5 hrs. ' After 4 hrs. 



■ An attempt was made to use potassium nitrate but it was noted that 
even in a neutral solution reduction of the nitrate to ammonia took place 
to an extent which could be noted with organic indicators. 

(9) 



The change in the hydrogen ion concentration was 
very rapid at first, but continually rose as long as the 
experiments were conducted. The plotted curves are 
all similar in character and of the general type shown 
in Fig. 5, which is for the yellow-gray silt loam with 
four tons of lime. 

Although equilibrium is not reached for a considerable 
period under the conditions of the experiment, the 
greater portion of the reaction takes place within a 
few minutes. After the neutral point is reached the 
hydrogen ion concentration continues to increase at 
the same relative rate without change. This phe- 
nomenon explains the condition noted by Maclntire' 
who distinguishes between immediate and continued 
lime requirement and Truog^ who makes the division 
into active and latent soil acidity. Both investigators 
added an excess of lime to the soil while in certain of 
the above experiments the lime requirements were not 
satisfied. Apparently the reaction is in each case an 
equilibrium reaction. Such being the case, any division 
as indicated by them is an arbitrary one, and "the 
results will not necessarily bear any relation to the 
total lime requirements. 





1 


1 1 1 1 1 


1 1 1 1 1- 


<l«> 


- 






«.w 


\ 


' 


- 


• r» 


\ 


. 




CM 


-"^ 


.,__ 


- 






- 


*.« 


1 


1 1 1 1 1 


1 1 1 1 1 



Fig. 5 

It would seem from an inspection of the curve that 

the relation may be expressed empirically by the 

equation 

dx 

- - K(A-.), 

in which x is the hydrogen ion concentration with 
time /, A the total change in hydrogen ion concentra- 

' Tenn. Expt. Sta., Bulletin 107 (1914), 193. 
> This Journal, 8 (1916), 341. 

(10) 



tion, and K a characteristic constant for the system 
under consideration. It would appear that the longer 
the interval chosen the more nearly would the results 
approach the maximum lime absorption, provided 
side reactions are not considered. It is quite probable 
that under field conditions the reactions will be far 
slower than the above experiment would indicate. The 
conditions in the field are often such that reactions 
cannot take place rapidly, while the converse is true 
in the laboratory. In this experiment a salt solution 
was used with calcium hydroxide, a condition, cer- 
tainly, where reaction may take place with utmost 
speed. 

This experiment would seem to throw some light 
upon the action of soil with potassium acetate in 
Expt. I. Potassium acetate is basic in character and 
even at the end of i8o hrs. the mixture with an 
acid soil contained only a slight excess of hydrogen 
ions. 

EXPERIMENT III. STUDY OF CHANGE IN HYDROGEN 
ION CONCENTRATION OF NEUTRAL SALT SOLUTION 
CONTAINING VARYING AMOUNTS OF BASE— 50 CC. of a 

0.5 N neutral salt solution containing the desired 
amount of base were placed in the gas cell and readings 
were taken at the end of 20 min. The gas cell was 
rocked continuously as in the former experiments. 

In the light of Expt. II, it can be readily under- 
stood that duplicates are somewhat difficult to obtain 
unless the manipulation is the same, and the time 
element eliminated. The speed of the reaction should 
depend somewhat upon the base used and the neutral 
salt with which it is combined. Considering these 
facts it is rather remarkable that the results obtained 
as shown in Table III should be of the same order. 
The experiments were all conducted upon yellow-gray 
silt loam. 

The salt solutions in contact with the soil (Series 
0.0 T. CaCOs) in Table III show considerable differ- 
ence in potential, but all show about the same neutral 
point, i. e., 5.0 T. CaCOa. From this point the variation 
in the readings becomes rather wide in both directions. 

The general form of the curves is shown by the type 
curve. Fig. 6. Two points regarding this curve may be 
discussed. It will be noted that the curve in the slope 
inclines more toward the vertical as the neutral point 
is reached and after this point is passed the slope 
inclines away from the vertical, approaching the slope 
at the lower end of the curve. 

(11) 



Table III — -Sti-dy ok Change in Hydrogen Ion Concentration op 
Neutral Salt Solution to Which Various Amounts op Base 
Have Been Added when Placed in Contact with Soil 



^Results given in volts. 



Soil is yellow-gray silt loam. 5 g were taken in 
each case) 



Base 


0.5 N 


0.5 N 


0.5 A^ 


0.5 iV 


5 A' 


Ivquivalent 


KCl 4- 


NaCI + 


CaClj + 


K2SO, + 


KCl + 


to T. CaCOa 


Ca(OH)j 


Ca(OH)2 


Ca(OH)2 


Ca(OH)i 


KOH 


0.0 


0.4843 


0.4995 


0.4966 


0.5245 


0.4843 


1.0 


0.5099 


0.5118 


0.5105 


0.5345 


0.5119 


2.0 


0.5247 


0.5180 


0.5259 


0.5527 


0.5269 


3.0 


0.5428 


0.5575 


0.5475 


0.5679 


0.5505 


4.0 


0.6019 


0.6125 


0.6215 


0.6169 


0.6136 


4.2 


0.6187 










4.4 


0.6342 










4.6 


. 6405 










4.8 


0.6651 










5.0 


0.6863 


o!6890 


0.6894 


o!676o 


0^6888 


5.2 


0.6943 










5.4 


0.7059 










5.6 


0.7211 










5.8 


0.7232 










6.0 


0.7387 


0.'74i4 


o!7670 


o!7330 


0^7391 


7.0 


0.7726 








0.7743 


8.0 


0.7942 










10.0 


0.8529 


o!8335 


o;7990 


o!8443 


0!88i9 



There is no abrupt change in hydrogen ion con- 
centration, as was noted by Hildebrand,* when a strong 
acid is neutralized by a base, but rather the curve 
approaches a straight line, i. e., for each addition of base 





1 1 1 1 1 


1 1 1 1 1 L 


HI 


- 




tn 


- 




«.rr 


- 


- 


<.» 


- 


^^-^-^'^^ 


:-!s 


- 


^^""^ 


».7I> 


- y^ 




its 


/ 




nu 


/ 




ISS 


^^,^ 


- 


».» 




- 


MS 


1 1 1 1 1 


1 1 I 1 1,1 



Tons CaCOj 

Fig. 6 

there is a corresponding nearly equal change in hydro- 
gen ion concentration. One may represent the change 
in the hydrogen ion concentration by the equation 

log C = KB -f Ki 
where C is the hydrogen ion concentration, B the base 
added, and K and Ki characteristic constants. 

The systems investigated above have all the charac- 
teristics of mixtures with a high reserve acidity, ^ i. e., 
the hydrogen ion concentration suffers very little 
change with comparatively large additions of base. 

' Loc. cil. 

' Washburn J. Am Chem Soc. 30 (1908), 37. 

(12) 



This would be expected if the acid is comparatively 
insoluble. 

Referring again to Table III it will be noted that 
after the neutral point is reached the hydrogen ion 
concentration for the system CaCl2 + Ca(0H)2 + 
Soil shows a higher hydrogen concentration than the 
system KCl + KOH + Soil for the same equivalents 
of base. This may be accounted for by the precipi- 
tating action of calcium hydroxide. The difference 
in conductivity adds weight to this view. 



»■!! 


1 


1 


1 




1 


1 


1 


1 1 


1- 


0.90 


- 


















tSS 


- 
















- 


30 


- 
















- 


a.rs 


- 
















- 


„. 


- 










^ 


/ 






OiS 


- 


















,« 


h- 






y^ 












Oif 






^^ 












~ 


,« 




' "^ 














- 


4f 


1 


1 


1 


1 


1 


1 


1 


1 


1 


iL 


3 / 


z 


3 


♦ 








8 f 





Toi\/s CaCO, 

Fig. 7 

When the base is allowed to act for a longer period 
of time there is in general a depression of the whole 
curve toward the acid side as is shown by Table IV 
and Fig. 7 which is the graph of the system 0.5 iV KCl 
+ Ca(0H)2 + Soil. 

Table IV — Effect op Time upon Absorption of Base 
(Soil used was yellow-gray silt loam. Time, 3 hrs. Results recorded in 





volts) 






0.5 N KCl -t- 


0.5 N KCl -1- 


T. CaC03 


Ca(OH)2 


KOH 


0.0 


0.4963 


0.4997 


1.0 


0.5120 


0.5157 


2.0 


0.5249 


0.5276 


3.0 


0.5418 


0.5445 


4.0 


0.5820 


0.5959 


5.0 


0.6397 


0.6488 


6.0 


0.6797 


0.6702 


7.0 


0.7217 


0.7268 



It will be noted that for the same amount of base 
there is a higher hydrogen ion concentration than 
shown in Table III. Hydrogen ion concentration 
increases with time in the presence of a soluble base. 

To compare the change of the hydrogen ion con- 
centration of solutions in contact with different soils, 
a number of soils were chosen and similar determina- 



tions made as with the yellow-gray silt loam in Table 
IV. All the soils were shaken for 3 hrs. with 50 cc of 
0.5 N potassium chloride solution containing the re- 
quired amount of lime. In each case 5 g. of soil 
were used, except in the case of the peat soils, of which 
2.5 g. were taken. The results are given in Table V. 

Graphs of some of the above readings (Fig. 8) show 
nearly straight line functions. Certain of the soils 
tested, Nos. i, 2. 3, 4, 6, and 11, show straight line 
functions for all measurements taken, while the others 
show more or less distinct changes in slope at two 
points, one being on the acid side, i. e., where the H"*" 
concentration is greater than io~^, and the other on 
the basic side. The same is true of the yellow-gray 
silt loam (Fig. 7). At no point in these graphs is there 
an abrupt change in slope, as would be expected if 
we were neutralizing a strong acid, but instead the 
neutral point is reached at an angle depending upon 
the character of the soil. It requires comparatively 
large additions of base to produce a marked change in 
hydrogen ion concentration. 

It may be readily understood that the organic 
indicators which change color at different hydrogen 
ion concentrations will show marked differences in the 
lime requirements of soils, because of the slight change 
in hydrogen ion concentration with the additions of 
lime. With the sandy loam soil a change in hydrogen 
ion concentration of the indicator from io~* (corre- 
sponding to the color change of methyl orange) to 
io~8 (corresponding to the color change of phenol- 
phthalein) would represent a change in the apparent 
lime requirement from 0.0 T. to 3 T.; and a change 
in the lime requirement for peat from about i T. to 
more than 20 T. of calcium carbonate. It is evident 
that soil solutions may be distinctly acid to litmus 
while alkaline to methyl orange. 

The slight changes in hydrogen ion concentration 
with addition of base will account for the great varia- 
tion in results obtained for the lime requirement of 
soils by the different methods proposed. The greater 
the slope of the curve the less will be the differences 
observed. With sandy soils it would be expected that 
the results obtained by various methods would approach 
each other, greater variations would be observed with 
other inorganic soils, while with soils high in organic 
matter it would be expected that the widest differences 
would be noted. 

(14) 



The time factor becomes important, as is shown in 
Tables I and II. The longer the base acts the higher 
is the hydrogen ion concentration at all initial concen- 
trations of base. The temperature undoubtedly should 
receive consideration, for the speed of reaction is 
increased, salts are hydrolyzed to a greater extent,^ 
and water of hydration may be decreased with increase 
in temperature. 



' 


'III 


1 1 1 1 1 1 U 


AfJ- 


- 


Shown Sandy LOiM 


Jf^ 


- 


^^/-^"^'^^ 


ots 


- 


^-■^'"'^ Chey PitSTic Cny~ 
^.^ ^ — ■""•rmovu.T lotM 


CSO 


^^ 


^ YiuoKmrYCiAY 


7e 


: /^ 


^^.^^^^^^^^^^^^^^ ' 


as 


- // ^ 


^y^^ aLKKCLWlOiS. =■ 


oto 


y ^ 


7^^^^^^~~ZZ^---—' — — ' — " 


0« 


^^Z^^^^---"^^ 


— 


O/o 


s^^"^ 


- 


Ui 


\ 1 i L_ 


1 1 1 1 1 1 : 



To MS CaCC^ 

Fig. 8 

We should expect that an acid soil would take up 
larger quantities of calcium from calcium carbonate 
than from calcium chloride, or sodium from sodium 
acetate than from sodium chloride as has been re- 
peatedly shown to be the case. This may be explained 
at least partially by differences in hydrolysis of the 
salts and the differences in the ionization of the acid. 
With a carbonate as the reagent, since the point of 
equilibrium depends also upon the partial pressure of 
carbon dioxide, it may be understood that the apparent 
lime requirement may be changed by simply changing 
the partial pressure of carbon dioxide. The results 
reported by Ames and Schellenberger,^ by the Hutchin- 
son- MacLennan^ and the vacuum methods, are in line 
with this argument. 

The placing of the soil under artificial conditions 
with respect to temperature, the use of powerful 
reagents and the extremely finely divided material 
employed in most of the methods suggested will 
certainly affect the speed of the reaction. Furthermore, 
the equilibrium is without doubt a metastable one. 

1 Landolt and Bomstein. "Physikalisch Chemische Tabellen." 

« This Journal, 8 (1916), 243. 

« Chem. News. 110 (1914), No. 2854, 61. 

(15) 



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(91) 



The present researches may not be said to be free from 
these faults. 

Table VI — Conductivity op Soil Solution Containing Various 

Quantities of Base, in Contact with Acid Soil 

(Base Calculated as Tons of CaCOs per Acre 

Results = Specific Conductivity X 10') 















Black 


Peatv 








Yellow-Gray Silt 


Yellow-Gray Silt 


Loam 


>,5> u,' 


•0 


Loam. Base 


Ca(OH) 


J Loam 


. Base 


KOH 


Base Ca(OH) . 


Ma a 

'^- 


•S^ 
















.T) (S 


U M 2 


■ 


bO 


Mi2 




bo 




■X bO 


UtJ 


■~ii --K 


5ii; 


.5 u 


.SK 


— U 


.S u 


B i2 


^ EO = 


SH 


•oK "O 
a noo 

tS— c« — 




nl — 


•o 

qoo 

(3 — 


_2K 






lac 
Loa 

Ca(' 
tandi 


C9 


'Ji C/3 


VI 


M 


Vi 


-/! 


m 


M 


n 7: 


0.0 


1.016 1.016 


0.913 


1.016 


1.016 


0.913 


2.341 


2.837 


3.813 


1.0 


1.089 1.129 


0.916 


1.276 


1.203 


1.132 


2.825 


3.388 


4 


015 


2.0 


1.220 1.129 


0.994 


1.733 


1.618 




3.050 


3.720 






3.0 


1.638 1.464 


1.146 


2.520 


2.376 


1.968 


3.468 


4.013 


4 


498 


4.0 


1.906 1.980 


1.495 


3.748 


3 . 245 


2.402 


3.744 


4.296 






4.4 


2.033 2.120 


1.691 


4.637 


3 . 748 












4.8 


2.226 2.248 


1.826 


5.100 


4.122 












5.2 


2.226 2.276 


1.980 


5.864 


4.620 












5.6 


2.346 2.365 


2.048 


7.318 


5.352 












6.0 


2.480 2.352 


2.074 


9.300 


5.978 


41888 


4!i40 


4^767 






7.0 


3.390 2.799 


2.310 


14.250 


8.718 




4.440 


4.968 


5 


408 


10.0 


10.170 4.842 


3.427 


43.580 


29.330 


25 .'420 


6.055 


5.457 


5 


820 



In soils in contact with solutions we have to deal 
with solid phases whose reaction velocities are neces- 
sarily very slow. The substances which produce the 
acid phenomena when salt solutions are added to acid 
soils are undoubtedly only slightly soluble and the 
products formed may pass into similar solid phases. 
Even substances which are assumed to be soluble may 
be held in the colloidal condition. Kahlenburg and 
Lincoln^ claim that the silicic acid of spring waters 
is present in the colloidal condition, and the fact that 
aluminum salts are hydrolyzed to a marked extent 
while the solubility of aluminum hydroxide is very 
slight may be taken as an indication that it is held in 
solution in part as a peptized colloid. The work of 
Mahin, Ingraham, and Stewart- seems to be evidence 
that the above statement is true in regard to the 
aluminates. 

In anticipation of criticism from the use of neutral 
salt solution in the potential readings throughout these 
investigations the following experiment was planned. 

It may be reasoned that if slightly soluble acids are 
present in acid soils it should be possible to measure 
the hydrogen ion concentration from such a water 
solution, if the acid ionizes at all, by means of the 
hydrogen electrode and other indicators. This has 
been done by Gillespie,^ but no attempts have been 
made to follow the change in hydrogen ion concentra- 

1 J. Phys. Chem., 2 (1898), 88. 

2 J. Am. Chem. Soc, 35 (1913), 30. 
' Loc. cit. 

(17) 



tion upon the addition of a base. In repeating Gil- 
lespie's experiment, using yellow-gray silt loam, it 
was found that after shaking 5 g. of the soil with 50 
cc. of distilled water for one-half hour, the potential 
reading of the mixture was depressed to 0.565 volt, 
but there is question whether equilibrium had yet been 
attained considering the fact that neutral salts depress 
the reading to about 0.5 volt. 

The above experiment was modified by using 50 cc. 
of distilled water containing the equivalent of 4 tons 
of lime calculated as calcium carbonate. After the 
readings had become constant 5 g. of yellow-gray silt 
loam were added and readings taken at stated intervals. 
The results are given in Table VII. 

Table VII — Change in Potential in an Acid Soil on Shaking with 
Solution of Calcium Hydroxide 
Time 
Minutes Volt 

0.9723 

30 0.7056 

60 0.6772 

90 0.6580 

240 0.6529 

420 0.6416 

Equilibrium evidently had not been reached after 7 
hrs. of continuous shaking. In performing the experi- 
ment some difficulty was encountered because of the 
high resistance of the chain, but the readings are typical. 
It will be noted that the reaction is much slower than 
in the presence of a neutral salt as may be expected 
from a consideration of the number of ions present in 
each case. 

The salt solutions used throughout this research, 
unless otherwise stated, were 0.5 iV, while the calcium 
hydroxide solution used above was 0.008 N, or an 
approximate ratio of 500 : 8 at the beginning of the 
reaction. While the neutral salt concentration re- 
mained practically unchanged the calcium hydroxide 
concentration, and consequently the hydroxyl ion 
concentration, became progressively less. 

The above experiment brings out another fact of 
importance. After the base has been neutralized the 
hydrogen ion concentration continues to rise slowly. 
The substance producing the action comes to an 
equilibrium very slowly with water, and 6 hrs. after 
the lime was neutralized hydrogen ions were still being 
thrown into the solution. 

EXPERIMENT IV. CONDUCTIVITY OF SOIL SOLU- 
TIONS — To test the question of the insolubility of the 
products formed when a base is added to an acid soil, 

(18) 



some conductivity experiments were conducted as a 
preliminary investigation. 

Five grams of soil were shaken with 75 cc. of distilled 
water containing the required amount of base and con- 
ductivity measurements made upon the recently 
shaken mixture of soil and solution with a plunge 
electrode having a constant of 0.0305. All readings 
were taken at 25° =»= 1°. 

From an inspection of Table VI, the following 
deductions may be made: 

The specific conductance increases with each addi- 
tion of base, but the increase is much greater with 
potassium hydroxide than with calcium hydroxide. 
The difference is far too great to be accounted for by 
the difference in conductivity of potassium and cal- 
cium ions.^ Either the calcium salts formed are less 
soluble or they must ionize to a far less extent than the 
potassium salts. The former view is substantiated 
by the difference noted in the soluble base left in solu- 
tion when a soil is treated with water containing an 
excess of base. From the standpoint of the absorption 
theory it may be argued that calcium hydroxide is 
absorbed to a greater extent than potassium hydroxide, 
but this theory can hardly be substantiated in face of 
the fact that potassium and calcium hydroxide seem 
to have practically equivalent power to neutralize the 
acid of the soil noted elsewhere in this paper. 

The specific conductivity decreases with time. This 
is true in every case except with the black peaty loam, 
where this may be accounted for by the increase in 
soluble material brought into solution. It will be 
noted that the conductivity of the water solution in 
contact with this soil increased after shaking, indi- 
cating that equilibrium had not been reached. If this 
is taken into account, the condvictivity of the solution 
in contact with this loam also decreases with time. 
It would seem that the reactions are progressive and 
that equilibrium is not reached for a considerable 
time period. This is directly in line with the change 
in hydrogen ion concentration determinations dis- 
cussed elsewhere in this paper. 

As the acid-producing substances present in the 
soil are extremely insoluble, as shown by the fact that 
the pure water extract shows little acidity with ordi- 

1 For method of calculating conductance for calcium and sodium ions 
see Bates, J. Am. Chem. Soc, 35 (1913), 534; and Washburn, "Principles 
of Physical Chemistry," New York, 1915, p. 214, for calculated values, 
the equivalent conductance of calcium and potassium ions at infinite dilu- 
tion being 51 and 63.3, respectively. 

(IQ) 



nary indicators, it would be expected that reactions 
would be exceedingly slow, especially near the equilib- 
rium point. It will be noted in this connection that 
there was a much more marked change in conduc- 
tivity with time at the higher concentrations of base. 

The rather slow increase in conductivity would 
point to precipitation effects which are more marked 
where calcium is used as the base. Acid substances 
which are highly insoluble would show high reserve 
acidity, a condition which is quite evident in the soils 
investigated. 

SUMMARY 

I — A new hydrogen electrode cell has been designed 
which has proved satisfactory for certain types of soil 
investigation. 

2 — When an acid soil is added to a neutral salt solu- 
tion the hydrogen ion concentration of the solution 
reaches a maximum almost immediately if the soil is 
wet thoroughly by the solution, but secondary reac- 
tions later cause a decrease in hydrogen ion concen- 
tration of the solution. 

3 — When an acid soil is added to a neutral salt solu- 
tion containing a free base the base is neutralized 
rapidly, as indicated by the change in hydrogen ion 
concentration of the solution, following closely the law 
for equilibrium reactions, but the hydrogen ion con- 
centration of the solution continues to rise for an 
unknown period. 

4 — There is no sharp break in the progress of base 
absorption by an acid soil which will warrant any 
arbitrary division, such as active and latent acidity, 
or immediate and eventual lime requirement. 

5 — The changes in log Ch approach nearly straight 
line functions with progressive addition of base in the 
presence of a neutral salt solution. 

6 — Different indicators will give differences in lime 
requirement for soils depending upon the slope of the 
log Ch curve. The greatest differences may be expected 
with soils high in organic matter. 

7 — A general explanation is given of the different 
results obtained by different methods for determining 
soil acidity. 

8 — When a base is added to an acid soil compara- 
tively insoluble products are formed. Calcium pro- 
duces a product less soluble than does potassium. 

9 — The specific conductance of a water solution of an 
(20) 



acid soil to which a base has been added increases with 
each addition of base, but the increase is greater with 
potassium than with calcium hydroxide, which is far 
too great to be accounted for by the difference in con- 
ductivity of potassium and calcium. 

lo — Calcium and potassium hydroxide have prac- 
tically equivalent power to neutralize the acid of an 
acid soil. 

II — The specific conductivity of a pure water solu- 
tion containing a base in contact with an acid soil 
decreases with time. 

1 2 — An acid soil shows high reserve acidity. 

13 — The reaction between a water solution of a base 
and an acid soil is much slower than in the presence of 
a neutral salt. 

14 — The absorption of bases by acid soils is due 
largely to relatively insoluble acids. 



(21) 



[Reprinted from the Journal of Industrial and Engineering Chemistry, 
Vol. 12, No. 6, page 559. June, 1920.] 



ACIDITY AND ACIDIMETRY OF SOILS. IH— COMPARI- 
SON OF METHODS FOR DETERMINING LIME RE- 
QUIREMENTS OF SOILS WITH HYDROGEN 
ELECTRODE. IV— PROPOSED METHOD 
FOR DETERMINATION OF LIME 
REQUIREMENTS OF SOILS 
By Henry G. Knight 

Oklahoma Agricultural and Mechanical College, Stillwater, 
Oklahoma 

Received October 14, 1919 
PART III 

Through the courtesy of Mr. J. W. Ames of the Ohio 
Agricultural Experiment Station fifteen samples of soil 
were obtained from a number of variously treated 
plots from one of the fertility sections of the Wooster 
farm located on silt loam soil which is derived from 
sandstones and shales. Ames and Schollenberger^ had 
made determinations of the lime requirement upon 
these soils by the Veitch, Hopkins, Hutchinson- Mac- 
Lennan, Maclntire, and vacuum methods and all were 
tested with litmus paper and found to give a decided 
reaction. "The west half of the plots had been treated 
with 1,875 lbs. per acre of calcium oxide in 1903, and 
2,000 lbs. of limestone in .1909. The composition of 
the lime materials applied was such that the equivalent 
of 5,700 lbs. of calcium carbonate had been applied to 
the limed halves of the plots previous to the time 
samples were taken from the plots, which was three 
years after the last treatment with lime."^ The 
amount of lime left was negligible. All samples gave 
a decided acid reaction to litmus and when mixtures 
of the soils with recently boiled distilled water were 
tested with the hydrogen electrode all showed a hydro- 
gen-ion concentration greater than io~^. 

Hydrogen-ion concentrations were determined upon 
these soils by use of the hydrogen electrode, using 
0.5 iV potassium chloride solution containing various 
predetermined amounts of lime. The samples were 
shaken for 3 hrs. preceding the readings given in Table 
I.' 

By interpolating the above results as straight line 
functions to determine the amount of lime necessary 

J This Journal, 8 (1916), 243. 

2 /did., 8 (1916), 224. 

• For method and apparatus see Part II, Ibid., 12 (1920), 45 7. 

(I) 



Tabub I — Hydrooen-Ion Concentrations of Solutions Containing 

Ohio Soils to Which Fixed Amounts of Lime Had Been Added 

(Readings are given in volts) 

T. IT. 2 T. 3 T. 4 T. 
Plot Fertilizer Lime Lime Lime Lime Lime 

None 0.5174 0.5684 0.6300 0.6739 0.7450 

None. Lime 0.5670 0.6406 0.6850 0.7393 

2 Acid Phosphate 0.5209 0.5725 0.6392 0.7000 

Acid Phosphate + Lime 0.5760 0.6561 0.6974 0.7473 

5 Sodium Nitrate 0.5184 0.5671 0.6317 0.6862 0.7482 

Sodium Nitrate + Lime. .. . 0.5803 0.6511 0.7103 0.7540 

I 1 Acid Phosphate + Mineral 

Potash + Sodium Nitrate 0.5164 0.5657 0.6308 0.6920 

Acid Phosphate 4- Mineral 
Potash + Sodium Nitrate 

+ Lime 0.5626 0.6398 0.6865 0.7421 

24 Acid Phosphate + Mineral 
Potash 4- Ammonium Sul- 
fate + Lime 0.5445 0.6230 0.6612 0.7260 

26 Bone Meal + Mineral Potash 

+ Sodium Nitrate 0.5209 0.5731 0.6375 0.6774 0.7394 

Bone Meal + Mineral Potash 

+ Sodium Nitrate + Lime 0.5516 0.6265 0.6841 0.7312 

29 Basic Slag + Mineral Potash 

4- Sodium Nitrate 0.5220 0.5794 0.6271 0.6826 0.7339 

Basic Slag + Mineral Potash 

+ Sodium Nitrate -f Lime . 5668 . 6424 . 7074 . 7467 

18 Manure 0.5202 0.5659 0.6164 0.6622 0.6938 

Manure + Lime 0.5574 0.6163 0.6708 0.7218 

to lower the hydrogen-ion concentration to io~^, 
using 0.69 volt as the potential at this concentration, 
a comparison may be made directly with the results 
given by Ames and Schollenberger.^ This comparison 
has been made in Table II. The results are given in 
pounds of lime as CaCOs required per acre. 

Table II — Comparison of the Amounts op Lime Required by Ohio 
Soils as Shown by Various Methods 

DO j,T3 

•| -9, £ '■= = 2o 

Plot Fertilizer K > S W > E 

1 None 3440 2000 3550 2925 7300 6456 

None, Lime 100 Alk. 2250 1700 4900 4225 

2 Acid Phosphate 2640 2000 3850 2700 7800 5670 

Acid Phosphate + Lime 80 Alk. 2400 975 3800 3640 

5 Sodium Nitrate 3640 1200 3550 2550 6200 6122 

Sodium Nitrate + Lime 120 Alk. 2500 1250 4225 3310 

1 1 Acid Phosphate + Mineral Potash 

+ Sodium Nitrate 3080 1800 3850 2825 7100 5960 

Acid Phosphate -f Mineral Potash 

+ Sodium Nitrate + Lime 80 Alk. 2500 1375 5900 4126 

24 Acid Phosphate -f Mineral Potash 

4- Ammonium Sulfate + Lime. . 4240 3000 4000 2700 8300 4889 
26 Bone Meal + Mineral Potash + 

Sodium Nitrate 2940 2000 3700 2250 7350 6407 

Bone Meal + Mineral Potash 4- 

Sodium Nitrate 4- Lime 360 Alk. 2900 1325 4050 4250 

29 Basic Slag 4- Mineral Potash 4- 

Sodium Nitrate 2560 1200 3600 2250 6600 6288 

Basic Slag 4- Mineral Potash 4- 

Sodium Nitrate 4- Lime 150 Alk. 2100 1075 4050 3460 

18 Manure 2760 2600 4200 3100 8500 7760 

Manure 4- Lime 120 Alk. 2950 1950 5200 4937 

'Ohio Experiment Station, Bulletin 306 (1916), 350. Hopkins' values 
are given as one-half that given in the table here, for which no explanation 
is offered. 

The following conclusions may be drawn from a 
study of Table II: 

' Ibid., 8 (1916), 244. 

(2) 



I — The vacuum method approaches nearer to the 
lime requirement as shown by the hydrogen electrode 
than do any of the other methods, but with this method 
the results are uniformly higher. It may be assumed 
that if the soils had been shaken with the lime for a 
longer period than 3 hrs. the lime requirement as shown 
by the hydrogen electrode would have approached 
that given by the vacuum method. 

2 — It is quite evident that the above methods, with 
the possible exception of the vacuum method, do not 
indicate the amount of lime necessary to comfiletely 
neutralize a soil, especially in the presence of neutral 
salts, except for a very limited period. 

PART IV 

The hydrogen electrode is not to be recommended 
for determination of the lime requirement of soils 
except as it may be valuable for checking other methods 
for the following reasons: 

I — It is difficult to manipulate even by one who has 
had considerable experience in using it. 

2 — The process is slower than any method so far 
proposed. 

3 — Expensive and delicate apparatus is required if 
satisfactory results are to be obtained. 

Although it may be found valuable for standardizing 
other methods, and approximate lime requirement 
values may be obtained by making two determinations 
and calculating the lime requirement as a linear func- 
tion, the method cannot be taken seriously as a com- 
mercial laboratory method. 

As the time element is an important factor auto- 
matic titration would not be entirely satisfactory. On 
account of the difficulties enumerated above, further 
investigations were made in the hope that a practical, 
rapid method for the determination of the lime require- 
ment of a soil could be worked out. 

Tacke's method^ would appear to have a logical 
foundation but in the light of the present investiga- 
tions it is difficult to conceive how it could be expected 
to yield concordant results, and it is doubtful if it or 
any of the proposed modifications would approach the 
actual amount of lime needed to bring the soil to true 
neutrality. 

An attempt was made to determine the lime require- 
ment by mixing a weighed quantity of soil with pre- 
cipitated calcium carbonate, adding recently boiled 
distilled water, boiling for a fixed period, and deter- 

» Chem..Ztg., 20 (1897), 174. 

(3) 



mining the evolved carbon dioxide by the Parr method^ 
as modified by Pettit.^ Concordant results could not 
be obtained, but by adding a neutral salt rather close 
results were 6btained. 

The method used for the results given in Tables I 
and II below was to take a weighed sample of soil, 
usually 5 or lo g., add an excess of precipitated calcium 
carbonate in a 125 cc. Erlenmeyer flask, attach to the 
Parr apparatus, run in about 25 cc. of normal salt 
solution,' and boil for a definite period.* The flask 

Table I7— Comparison of Lime Requirements of Ohio Soils as Show.v 

Bv Various Methods, Including the Proposed Modified Tacke 

Method, AND the Differences Shown by Limed and Unlimed 

Plots. Previous Application op Lime on Limed Plots 

5700 Lbs. per Acre 



•a — 



3 ot: ua.- U3 V 



Plot Fertilizer ffi>2K>a 22 

None 3440 2000 3550 2925 7300 6456 6941 7783 

None + Lime 100 Alk. 2250 1700 4900 4225 6908 

Difference 3340 .. 1300 1225 2400 2231 .. 875 

2 Acid Phosphate 2640 2000 3850 2700 7800 5670 6737 7809 

Acid Phosphate + Lime 80 Alk. 2400 975 3800 3640 5123 6599 

Difference 2560 .. 1450 1725 4000 2030 1644 1210 

5 Sodium Nitrate 3640 1200 3550 2550 6200 6122 7264 7924 

Sodium Nitrate + Lime. 120 Alk. 2500 1250 4225 3310 5431 

Difference 3520 .. 1050 1300 1975 2812 .. 2493 

1 1 Acid Phosphate + Min- 
eral Potash + Sodium 

Nitrate 3080 1800 3850 2825 7100 5960 7042 7493 

Acid Phosphate + Min- 
eral Potash + Sodium 

Nitrate + Lime 80 Alk. 2500 1375 5900 4126 .. 6184 

Difference 3000 1350 1450 1200 1834 .. 1309 

24 Acid Phospnate + Min- 
eral Potash + Sodium 
Nitrate -f- Ammonium 

Sulfate -I- Lime 4240 3000 4000 2700 8300 4889 . . 6547 

26 Bone Meal -|- Mineral 
Potash -f- Sodium Ni- 
trate 2940 2000 3700 2250 7350 6407 6973 7530 

Bone Meal -|- Mineral 
Potash 4- Sodium Ni- 
trate -f- Lime 360 Alk. 2900 1325 4050 4250 6951 

Difference 2580 800 925 3300 2157 579 

29 Basic Slag -|- Mineral 
Potash -f Sodium Ni- 
trate 2560 1200 3600 2250 6600 6288 . . 7822 

Basic Slag -f- Mineral 
Potash -+- Sodium Ni- 
trate + Lime 150 Alk. 2100 1075 4060 3460 . . . 6233 

Difference 2410 .. 1500 1175 2540 2828 1589 

18 Manure 2760 2600 4200 3100 8500 7760 8327 8776 

Manure -f Lime 120 Alk. 2950 1950 5200 4937 7206 

Difference 2640 1250 1150 3300 2823 .. 1570 

> J. Am. Chem. Soc, 26 (1904). 294. 

» Ibid., 1640. 

» Preliminary experiments showed that it made little difference what 
neutral salt was used. Experiments were made with definite quantities 
of KCl, NaCl and KNOi; also by varying the amount of calcium carbonate 
and KCl with duplicate results, provided there was an excess of calcium 
carbonate and not enough neutral salt to change materially the boiling 
point of the solution. 

* If the boiling was attended with frothing a few drops of a neutral oil 
were added. 

(4) 



and condenser were filled with distilled water to a 
mark upon the capillary tube connecting the con- 
denser with the eudiometer. Readings were taken as 
instructed by Pettit.^ It is recommended that lo 
min. be fixed as the time for boiling. 

The soils given in Table I are the ones described by 
Ames and SchoUenberger- and received by the writer 
from them. The results are calculated in pounds of 
calcium carbonate to 2,000,000 lbs. of soil, except in 
case of peat soils where 1,000,000 lbs. were used as a 
basis. 

Other soils investigated are given in Table II. 
Hopkins and hydrogen electrode determinations were 
made and the proposed modified Tacke determinations 
were made after boiling for 5, 10, and 30 min. 

Table II — Lime Requirements op Soils as Shown by the Hopkins, 

THE Hydrogen Electrode, and the Proposed Modified 

Tacke Methods' 

Proposed Modified 
Tacke Method 

Hydrogen 5 10 30 

Soils Hopkins Electrode Min. Min. Min. 

Yellow-Gray Silt Loam 8400 12490 14847 15822 

Black Peaty Loam 6890 32420 

Black Clay Loam 9964 38026 ... 34718 

Peat 6535 34260 1101 1 20772 

Black Muck 8112 («) ... 31342 

Yellow Silt Clay 10160 14708 

Gray Clayey Silt 15840 («) 28803 30640 33624 

Gray Plastic Clay 5200 10000 ... 1 1094 

Yellow Plastic Clayey Silt 9340 14650 18173 19611 

Yellow Silt Loam 5780 14578 13453 14954 17250 

Brown Sandy Loam 3300 4574 ... 5792 

' For descriptions of soils see Part I, This Journal, 12, (1920), 340. 

2 More than 40000. 

» More than 20000. 



It will be noted that the proposed modification of 
Tacke's method gives varying results which depend 
upon the time of boiling, and in every case except for 
peat and yellow silt loam, Table II, the s-min. 
boiling period showed a higher lime requirement than 
the hydrogen electrode. A lo-min. boiling period 
showed a higher lime requirement in every case except 
one (Ohio Plot No. 24, lime, Table I) than did the 
vacuum method proposed by Ames and SchoUen- 
berger.^ That the reaction is not complete even at the 
end of the lo-min. period is shown by the increase at 
the end of a 30-min. boiling period. 

The proposed method has the advantage over most 
of the others in that it gives a figure which represents 
the "power of a soil to decompose calcium carbonate" 

' hoc. cit. 

» This Journal, 8 (1916), 243. 

» Loc. cit.. Part III. 

(5) 



(which may be assumed to be a measure of the eventual 
lime absorption), is rapid, and approximates the re- 
sults obtained with the hydrogen electrode. 

It is quite apparent that the interaction at room 
temperature between a soil and a caustic lime solution 
containing a neutral salt is not complete within 3 hrs. 
(which is the period at which hydrogen-ion concentra- 
tions were determined with the hydrogen electrode), 
as shown by both the vacuum and the modified Tacke 
methods. The period required for the completion of 
the interaction at room temperature between a soil 
and calcium carbonate may be assumed to be con- 
siderable for Maclntire^ has shown that calcium passes 
slowly into the form of silicates. 

It is quite evident that any of the proposed methods 
gives comparative results only. The true lime hunger 
as it relates to cropping is after all the matter that we 
are most interested in determining, and this, it seems, 
with the present state of our knowledge must be deter- 
mined by field experiments. 

MEASUREMENT OF THE REDUCTION OF ACIDITY 

By subtracting the acidity values found in acid soils 
from those found in unlimed soils from the correspond- 
ing half plots at the Ohio Station (Table I) differences 
are obtained which measure the residual reduction in 
acidity due to previous applications equivalent to 5,700 
lbs. of calcium carbonate. As an average of the results 
from seven plots which afford data for this measure- 
ment the reduction in acidity is 2,864 lbs. by the 
Hopkins method, 2,674 by the vacuum method, and 
2,388 by the hydrogen electrode, while the Maclntire, 
Hutchinson and modified Tacke methods show reduc- 
tions of 1,243, 1)279, and 1,375 lbs., respectively. That 
any method will show a greater reduction in acidity 
than actually occurs and remains at the time of sam- 
pling seems extremely doubtful. 

As suggested above, the vacuum method appears to 
furnish the most trustworthy measure of the total lime 
requirement and it also seems safe to assume that the 
hydrogen electrode will give results in substantial 
agreement with the vacuum method if sufficient time 
is allowed. If these methods are accepted as standards, 
then the Hopkins method seems to give correct results 
when used to measure the reduction in soil acidity by 
applications of lime. It may also measure with accu- 
racy the most immediate lime need, although it does not 

' Tennessee Experiment Station, Bulletin 107 (1914). 
^6) 



measure the total power of a soil to decompose car- 
bonates. 

If we assume that the reduction in acidity should 
be approximately the same for all limed plots, the 
Hopkins method and the hydrogen electrode show 
the highest percentage consistency. 

CONCLUSIONS 

I — A method has been proposed for determining 
the power of a soil to decompose calcium carbonate 
which approximates the results obtained by use of the 
hydrogen electrode. 

2 — The Hopkins and the hydrogen electrode methods 
show the highest percentage consistency for measuring 
the reduction of acidity for limed soils. 



(7) 



VITA 



Henry Granger Knight was born at Bennington, 
Ottawa Co., Kansas, July 21, 1878. He secured his 
common school education in the public schools of 
Kansas and Washington. The first two years of high 
school work was pursued at Leland, Washington- the 
third year was completed at the Port Townsend High 
School, from which institution he received a diploma in 
1896. In the fall of 1897 he entered the University of 
Washington at Seattle as a conditioned freshman. He 
graduated from this institution in June 1902 with the 
ciegree ot Bachelor ot Arts with honors in Chemistry. 
During the years 1898-1899, 1900, 1901-1902, he held 
the position of Assistant in Chemistry at the Univer- 
sity of Washington. 

He entered the University of Chicago in July, 
1902, as a member of the Graduate School where he 
studied chemistry under the direction of Dr. McCoy 
until August 1903, at which time he left to accept a 
position of Assistant Professor of Chemistry at the 
University of Washington. In January 1904 he was 
elected Professor of Chemistry at the University of 
Wyoming. He was granted the degree of Master of 
Arts by the University of Washington in 1904 upon 
work at the University of Washington, the University 
of Chicago and a thesis completed at the University of 
Wyoming. While holding this position in the summer 
of 1906 he attended the summer school of the Univer- 
sity of Chicago. In June 1910, he was appointed 
Director of the University of Wyoming Agricultural 
Experiment Station and Agricultural Chemist of the 
Agricultural College and one year later was also ap- 
pointed Dean of the Agricultural College. In Septem- 
ber 1915, he was granted a leave of absence to carry on 
graduate work at the University of Illinois. 



He is the senior author of Bulletin No. 62, "Some 
F'ood Products and their Adulteration;" Bulletin No. 
65, "Wyoming Forage Plants and their Chemical Com- 
position — Studies No. 1;" Bulletin No. 69, "Digestion 
Experiments with Wethers;" Bulletin No. 70, "Wyo- 
ming Forage Plants and their Chemical Composition — 
Studies No. 2;" Bulletin No. 76, "Wyoming Forage 
Plants and their Chemical Composition — Studies No. 
3;" Bulletin No. 78, "Digestion Kxperiments II;" 
Bulletin No. 82, "Soil Nitrogen;" Bulletin No. 87, 
"Wyoming Forage Plants and their Chemical Compo- 
sition — Studies No. 4; " Bulletin No. 97, "The Identifi- 
cation of the Woody Aster;" Bulletin No. 100, 
"Meteorology for Twenty Years;" and "Alkali VI," 
Sixteenth Annual Report; all of which have been pub- 
lished by the University of WVoming Agricultural Ex- 
periment Station. He is the junior author of "Notes 
on Qualitative Analysis," published by John Wiley and 
Sons, and Bulletin No. 94, "The Chemical Examination 
of Death Camas," published by the Univcrsitv of 
Wyoming Agricultural Experiment Station. 

In 1903 he was Fellow elect at Chicago. He is a 
member of the W'ashington Chapter of Sigma Xi and 
of Phi Beta Kappa, and of the Illinois Chapter of 
Sigma Upsilon, The American Chemical Society, 
National Geographic Society, the Society for the Pro- 
motion of Agricultural Science, et cetera. 



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